 In this video, we are going to introduce ourselves to a new class of elements, the D-Block elements. The D-Block, as we can see here, extends from group 3 all the way to group 12. They are an integral part of chemistry and can be credited with a number of interesting innovations of the past century, be it in engineering, catalysis or biomedical fields. For example, iron and its alloys are widely used in construction and manufacturing industries. Titanium is used in dental implants and also in artificial hip replacement surgeries. Manganese, another D-Block element is widely used in dry battery cells. Similarly, you can find many such applications of the D-Block elements in our everyday lives. But to understand what makes these elements so versatile in their applications, it is important to study their chemistry. And in this particular video, we will take a brief look at the position of the D-Block elements and also look at some interesting nuances about their electronic configurations. As I said before, the D-Block spans from group 3 to group 12. And it comprises of four series of elements that are formed by the filling of 3D, 4D, 5D and 60 subtiles of electrons. These elements are also called transition metals. Because firstly, they are all metals. And secondly, their properties are kind of transitional between the highly reactive S-Block elements that predominantly form ionic compounds and that of the P-Block elements that predominantly form covalent compounds. Now, depending on the oxidation state of the transition metals, the type of bonding in these metals can range from ionic to covalent bonding. Now, another interesting thing is that not all of the D-Block elements can be referred to as transition metals. There are a few exceptions which we will discuss shortly. But for now, let's talk about their electronic configurations. All of us remember the Off-Bow principle, right? Off-Bow principle states that the filling of the orbitals takes place based on the energy of the orbitals. Electrons first fill the subtiles of the lowest energy and then go on to fill the higher energy orbitals. Now, based on this, we get a certain order as we can see here. So, we know that in the first series of the D-Block elements, the filling of the 3D orbitals take place, right? But if you compare 3D with forest subtiles, which one is higher in energy? The one that has the higher value of N plus L value will have the greater energy, right? Now, N plus L value of 3D subtiles is 5, whereas that of 4S is 4. That is, 3D is greater in energy than 4S. And therefore, obviously, the 4S subtiles will get filled first, followed by 3D. For example, if you look at potassium, which has 19 electrons, the electronic configuration would be argon 4S1 and not 3D1 because the filling of 4S subtiles takes place first and not 3D because 4S is lower in energy. Similarly, calcium, which has 20 electrons, has an electronic configuration of argon 4S2 and not 3D2. Now, the element that follows calcium is a D-Block element called scandium, which has 21 electrons. But strangely, something interesting happens between calcium and scandium. Between calcium and scandium, this order of the energy level kind of reverses. 4S becomes slightly higher in energy than 3D. But that's not the only weird part. You see, with 21 electrons, if 3D becomes lower in energy than 4S, does it mean that the electronic configuration of scandium would become argon 3D3? You know, because it has 3 electrons and instead of 4S, do all the 3 electrons occupy the 3D orbital? Actually, no. Well, it turns out that because 3D has slightly less energy than 4S, the first electron of course enters the 3D orbital, but the second and the third electrons enter the 4S orbital. So that means the actual electronic configuration of scandium would be argon 3D1 4S2 and this is the correct configuration of scandium. I know this looks super weird, you know, according to whatever we learned about the electronic configuration and the filling of orbitals, but this random arrangement or distribution of electrons takes place due to a combination of several complex factors. You see, although we study the order of energy levels as shown here, you know, as if they are fixed energy levels, in reality the orbital energy is not fixed per se. The energy of the orbitals depend on several factors like electron filling, mutual electron-electron repulsion, nuclear charge and many other factors. And especially when the two levels are close in energy like 4S and 3D, there is a higher chance of, how do I say, the electrons jumping between these energy levels. Now, why does the 2 electron go into the higher 4S sub-shell instead of continuing to fill up the 3D orbitals? Well, unfortunately, we cannot get into the details here, but turns out that even though it is occupying the higher energy 4S orbitals here, overall this might be a more stabilizing state for the atom as such. And this is why instead of continuing in the 3D orbital, the second and the third electron fills up the 4S sub-shell and after that the subsequent electrons go ahead and fill the rest of the 3D orbitals. Nevertheless, for now, all you need to remember is that until calcium, like how we have learned in Off-Boss principle, 4S sub-shell is lower in energy than 3D, but from scandium, it so happens that this order kind of reverses and 4S becomes slightly higher in energy than 3D. And this is why the electron first fills the 3D orbital and the subsequent electrons goes into 4S and so on. Now, let's look at titanium, the next element which has 22 electrons. Now, titanium, the electronic configuration would be argon. What do you think it would be? First, the 3D levels get filled, right? The first electron goes into 3D orbital. The second and the third electrons go into 4S orbital. And the fourth electron, of course, because 4S cannot accommodate any more electrons here, it goes back to the 3D orbital. So, the electronic configuration of titanium would be argon, 3D2, 4S2. Similarly, vanadium, which has 23 electrons, the electronic configuration would be argon, 3D3, 4S2 and so on. Now, here's a question. While forming ions, from which orbital do you think the electrons would be lost? For example, if you have vanadium 2+, which means you have 21 electrons, from which of the orbitals does the electron leave? Does it go from the 4S orbital or does it go from one of the 3D orbitals? Pause the video and think about it for a moment, okay? Well, we know that electrons are always ejected or removed from the higher energy level, right? And here, because the 4S level is slightly higher in energy than the 3D, this is why the electrons are lost first from the 4S orbitals and then from the 3D orbitals. That means vanadium 2+, will have the following electronic configuration, and not 3D1, 4S2. This is correct and this is wrong. That is, the electrons are lost from the 4S orbitals and not from the 3D orbital because 4S is slightly higher in energy than the 3D. So, this has been confirmed through anisation studies that showed that the electron was actually ejected or removed from the 4S orbital and not from 3D. So, we saw scandium, titanium, vanadium. The next element of the D-block is chromium. Chromium has a total of 24 electrons and that means based on what we just studied, its electronic configuration should be argon 3D4, 4S2, right? But this is not what we actually see experimentally. You see, here the electron filling takes place in such a way that we get a half filled electronic configuration. Yes, that is, the electronic configuration would be argon 3D5, 4S1, half filled D orbitals. Now, this has been argued on the basis that half filled electronic configuration offers extra stability. From manganese again, we see similar filling of the electrons until we again reach copper. Copper has 29 electrons and here again, unlike what we expect that the electronic configuration which is argon 3D9, 4S2, the actual act configuration is a completely filled 3D sub-shell and half filled 4S sub-shell. So, this is what is the actual configuration of copper which is argon 3D10, 4S1. Here again, the reasoning is based on the stability of the fully filled D orbitals. Now, remember I mentioned that not all D-block elements are transition metals? Well, that's because according to the latest IUPAC definitions the transition metals are those which have incomplete D-sub-shell either in neutral atom or in their ions. And because zinc, cadmium and mercury have completely filled D orbitals, we don't generally consider them transition metals. The common oxidation state here is actually plus 2 which means even after losing two electrons in the S orbital they still have completely filled D orbitals. In other words, their properties are slightly different from the other elements of the series. So, even though zinc, cadmium and mercury are excluded from the transition series elements, they are still D-block elements, right? And therefore, we study their chemistry along with the rest of the D-block elements.